NCERT Solutions for Class 10 Science Chapter 1 in English: Chemical Reactions and Equations (In-Text & Exercises)

Answer: Magnesium is a highly reactive metal. When exposed to atmospheric air, it slowly reacts with oxygen to form a thin, stable, and inert layer of magnesium oxide (MgO) on its surface.

This white protective layer prevents further contact between the fresh magnesium metal and atmospheric oxygen, hindering burning. Therefore, before burning in air, the magnesium ribbon is thoroughly rubbed and cleaned with sandpaper to remove this oxide layer, allowing the metal to ignite readily and burn with a dazzling white flame.

Chemical Equation:
$$2\text{Mg (s)} + \text{O}_2\text{ (g)} \xrightarrow{\Delta} 2\text{MgO (s)}$$

Answer: (i) Hydrogen and Chlorine:
$$\text{H}_2\text{ (g)} + \text{Cl}_2\text{ (g)} \longrightarrow 2\text{HCl (g)}$$
(ii) Barium chloride and Aluminium sulphate:
$$3\text{BaCl}_2\text{ (aq)} + \text{Al}_2(\text{SO}_4)_3\text{ (aq)} \longrightarrow 3\text{BaSO}_4\text{ (s)} \downarrow + 2\text{AlCl}_3\text{ (aq)}$$
(White precipitate of barium sulphate is formed.)

(iii) Sodium and Water:
$$2\text{Na (s)} + 2\text{H}_2\text{O (l)} \longrightarrow 2\text{NaOH (aq)} + \text{H}_2\text{ (g)} \uparrow + \text{Heat}$$

Answer: (i) Precipitation reaction between Barium chloride and Sodium sulphate:
$$\text{BaCl}_2\text{ (aq)} + \text{Na}_2\text{SO}_4\text{ (aq)} \longrightarrow \text{BaSO}_4\text{ (s)} \downarrow + 2\text{NaCl (aq)}$$
(State symbols: (aq) = aqueous solution, (s)↓ = solid insoluble precipitate)

(ii) Neutralisation reaction (Sodium hydroxide + Hydrochloric acid):
$$\text{NaOH (aq)} + \text{HCl (aq)} \longrightarrow \text{NaCl (aq)} + \text{H}_2\text{O (l)}$$

Answer: (i) Name and Chemical Formula of Substance ‘X’:
The substance ‘X’ is Calcium Oxide, commonly known as Quicklime.
Its chemical formula is $$\text{CaO}$$.

(ii) Reaction of Substance ‘X’ with Water:
Calcium oxide reacts vigorously with water to produce slaked lime (calcium hydroxide) with the liberation of a huge amount of heat (exothermic combination reaction):
$$\text{CaO (s)} + \text{H}_2\text{O (l)} \longrightarrow \text{Ca(OH)}_2\text{ (aq)} + \text{Heat}$$

Answer: Water is composed of two hydrogen atoms combined with one oxygen atom by volume, as represented by its chemical formula $$\text{H}_2\text{O}$$ ($2:1$ ratio).

During the electrolytic decomposition of water, an electric current breaks water molecules into hydrogen gas and oxygen gas according to the balanced stoichiometric equation:
$$2\text{H}_2\text{O (l)} \xrightarrow{\text{Electricity}} 2\text{H}_2\text{ (g) [at Cathode]} + \text{O}_2\text{ (g) [at Anode]}$$
As observed, 2 molecules of water yield 2 volumes of hydrogen gas at the negative cathode and only 1 volume of oxygen gas at the positive anode. Thus, the ratio of hydrogen to oxygen is $2:1$.
Name of the gas with double volume: Hydrogen ($\text{H}_2$).

Answer: In the reactivity series of metals, Iron (Fe) is more reactive than Copper (Cu).

When an iron nail is dipped into a blue solution of copper sulphate, iron displaces copper from its solution to form ferrous sulphate (Iron(II) sulphate):
$$\text{Fe (s)} + \text{CuSO}_4\text{ (aq) [Blue]} \longrightarrow \text{FeSO}_4\text{ (aq) [Pale Green]} + \text{Cu (s) [Brown coating]}$$
Due to the consumption of blue $\text{Cu}^{2+}$ ions and the formation of light green $\text{Fe}^{2+}$ ions in the solution, the blue colour fades to pale green, and reddish-brown copper metal deposits onto the surface of the iron nail.

Answer: A prime example is the precipitation reaction between aqueous solutions of Lead(II) nitrate and Potassium iodide:

$$\text{Pb(NO}_3)_2\text{ (aq)} + 2\text{KI (aq)} \longrightarrow \text{PbI}_2\text{ (s)} \downarrow \text{ [Yellow Precipitate]} + 2\text{KNO}_3\text{ (aq)}$$
In this reaction, the lead ions ($\text{Pb}^{2+}$) and potassium ions ($\text{K}^+$) mutually exchange their anions ($\text{NO}_3^-$ and $\text{I}^-$) to form a brilliant yellow precipitate of lead iodide.

Answer: (i) In reaction $4\text{Na} + \text{O}_2 \longrightarrow 2\text{Na}_2\text{O}$:
• Substance Oxidised: Sodium ($\text{Na}$) — because it gains oxygen to form $\text{Na}_2\text{O}$.
• Substance Reduced: Oxygen ($\text{O}_2$) — because it gains electrons.

(ii) In reaction $\text{CuO} + \text{H}_2 \longrightarrow \text{Cu} + \text{H}_2\text{O}$:
• Substance Oxidised: Hydrogen ($\text{H}_2$) — because it gains oxygen to form water ($\text{H}_2\text{O}$).
• Substance Reduced: Copper(II) oxide ($\text{CuO}$) — because it loses oxygen to form metallic copper ($\text{Cu}$).

Correct Answer: (i) (a) and (b) are incorrect.

Scientific Explanation:
In this redox reaction:
• Lead oxide ($\text{PbO}$) loses oxygen to form lead ($\text{Pb}$). Thus, lead oxide is reduced (statement d is correct, whereas statement a is incorrect).
• Carbon ($\text{C}$) gains oxygen to form carbon dioxide ($\text{CO}_2$). Thus, carbon is oxidised (statement c is correct, whereas statement b is incorrect).
Hence, statements (a) and (b) are the incorrect ones.

Correct Answer: (d) displacement reaction.

Scientific Explanation:
Aluminium ($\text{Al}$) is more reactive than iron ($\text{Fe}$). In this reaction, aluminium displaces iron from iron(III) oxide to form aluminium oxide and molten iron. This high-temperature reaction is also known as the thermite reaction.

Correct Answer: (a) Hydrogen gas and iron chloride are produced.

Chemical Reaction:
$$\text{Fe (s)} + 2\text{HCl (aq)} \longrightarrow \text{FeCl}_2\text{ (aq)} + \text{H}_2\text{ (g)} \uparrow$$
Iron displaces hydrogen from dilute hydrochloric acid, forming iron(II) chloride and releasing hydrogen gas with effervescence.

Definition of Balanced Chemical Equation:
A chemical equation in which the total number of atoms of each element is equal on both the reactants side and the products side is called a balanced chemical equation.

Why Balancing is Necessary:
Chemical equations must be balanced to satisfy the Law of Conservation of Mass formulated by Antoine Lavoisier.
According to this law: “Mass can neither be created nor destroyed in a chemical reaction.”
Therefore, the total mass of the products must equal the total mass of the reactants. This condition is only satisfied when the number of atoms of every element remains conserved before and after the reaction.

(a) Nitrogen + Hydrogen → Ammonia:
$$\text{N}_2\text{ (g)} + 3\text{H}_2\text{ (g)} \longrightarrow 2\text{NH}_3\text{ (g)}$$

(b) Burning of Hydrogen sulphide in air:
$$2\text{H}_2\text{S (g)} + 3\text{O}_2\text{ (g)} \longrightarrow 2\text{H}_2\text{O (l)} + 2\text{SO}_2\text{ (g)}$$

(c) Barium chloride + Aluminium sulphate:
$$3\text{BaCl}_2\text{ (aq)} + \text{Al}_2(\text{SO}_4)_3\text{ (aq)} \longrightarrow 2\text{AlCl}_3\text{ (aq)} + 3\text{BaSO}_4\text{ (s)} \downarrow$$

(d) Potassium metal + Water:
$$2\text{K (s)} + 2\text{H}_2\text{O (l)} \longrightarrow 2\text{KOH (aq)} + \text{H}_2\text{ (g)} \uparrow + \text{Heat}$$

(a) Nitric acid + Calcium hydroxide:
$$2\text{HNO}_3\text{ (aq)} + \text{Ca(OH)}_2\text{ (aq)} \longrightarrow \text{Ca(NO}_3)_2\text{ (aq)} + 2\text{H}_2\text{O (l)}$$

(b) Sodium hydroxide + Sulphuric acid:
$$2\text{NaOH (aq)} + \text{H}_2\text{SO}_4\text{ (aq)} \longrightarrow \text{Na}_2\text{SO}_4\text{ (aq)} + 2\text{H}_2\text{O (l)}$$

(c) Sodium chloride + Silver nitrate:
$$\text{NaCl (aq)} + \text{AgNO}_3\text{ (aq)} \longrightarrow \text{AgCl (s)} \downarrow + \text{NaNO}_3\text{ (aq)}$$
(Already balanced with 1:1:1:1 stoichiometry.)

(d) Barium chloride + Sulphuric acid:
$$\text{BaCl}_2\text{ (aq)} + \text{H}_2\text{SO}_4\text{ (aq)} \longrightarrow \text{BaSO}_4\text{ (s)} \downarrow + 2\text{HCl (aq)}$$

(a) Calcium hydroxide + Carbon dioxide:
$$\text{Ca(OH)}_2\text{ (aq)} + \text{CO}_2\text{ (g)} \longrightarrow \text{CaCO}_3\text{ (s)} \downarrow + \text{H}_2\text{O (l)}$$

(b) Zinc + Silver nitrate:
$$\text{Zn (s)} + 2\text{AgNO}_3\text{ (aq)} \longrightarrow \text{Zn(NO}_3)_2\text{ (aq)} + 2\text{Ag (s)}$$

(c) Aluminium + Copper(II) chloride:
$$2\text{Al (s)} + 3\text{CuCl}_2\text{ (aq)} \longrightarrow 2\text{AlCl}_3\text{ (aq)} + 3\text{Cu (s)}$$

(d) Barium chloride + Potassium sulphate:
$$\text{BaCl}_2\text{ (aq)} + \text{K}_2\text{SO}_4\text{ (aq)} \longrightarrow \text{BaSO}_4\text{ (s)} \downarrow + 2\text{KCl (aq)}$$

(a) Equation:
$$2\text{KBr (aq)} + \text{BaI}_2\text{ (aq)} \longrightarrow 2\text{KI (aq)} + \text{BaBr}_2\text{ (s)}$$
• Reaction Type: Double displacement and precipitation reaction

(b) Equation:
$$\text{ZnCO}_3\text{ (s)} \xrightarrow{\Delta} \text{ZnO (s)} + \text{CO}_2\text{ (g)}$$
• Reaction Type: Thermal decomposition reaction

(c) Equation:
$$\text{H}_2\text{ (g)} + \text{Cl}_2\text{ (g)} \longrightarrow 2\text{HCl (g)}$$
• Reaction Type: Combination reaction

(d) Equation:
$$\text{Mg (s)} + 2\text{HCl (aq)} \longrightarrow \text{MgCl}_2\text{ (aq)} + \text{H}_2\text{ (g)} \uparrow$$
• Reaction Type: Displacement reaction

1. Exothermic Reactions:
Reactions in which energy is released in the form of heat along with the formation of products are called exothermic reactions.
• Example 1 (Combustion of methane):
$$\text{CH}_4\text{ (g)} + 2\text{O}_2\text{ (g)} \longrightarrow \text{CO}_2\text{ (g)} + 2\text{H}_2\text{O (g)} + \text{Heat}$$
• Example 2 (Reaction of quicklime with water):
$$\text{CaO (s)} + \text{H}_2\text{O (l)} \longrightarrow \text{Ca(OH)}_2\text{ (aq)} + \text{Heat}$$

2. Endothermic Reactions:
Reactions in which energy is absorbed from the surroundings in the form of heat, light, or electricity to break bonds in reactants are called endothermic reactions.
• Example 1 (Thermal decomposition of limestone):
$$\text{CaCO}_3\text{ (s)} \xrightarrow{\text{Heat (}\Delta\text{)}} \text{CaO (s)} + \text{CO}_2\text{ (g)}$$
• Example 2 (Photolytic decomposition of silver chloride):
$$2\text{AgCl (s)} \xrightarrow{\text{Sunlight}} 2\text{Ag (s)} + \text{Cl}_2\text{ (g)}$$

To stay alive, our body requires energy, which is derived from the food we ingest.

During digestion, large complex food items are broken down into simpler substances. For example, carbohydrate-rich foods like rice, potatoes, and bread are digested into glucose ($\text{C}_6\text{H}_{12}\text{O}_6$).

During respiration, oxygen inhaled from the air enters our body cells and oxidises glucose, releasing carbon dioxide, water, and substantial energy in the form of ATP molecules:
$$\text{C}_6\text{H}_{12}\text{O}_6\text{ (aq)} + 6\text{O}_2\text{ (aq)} \longrightarrow 6\text{CO}_2\text{ (aq)} + 6\text{H}_2\text{O (l)} + \text{Energy (ATP)}$$
Since metabolic energy/heat is liberated during this oxidation process, cellular respiration is classified as an exothermic reaction.

Decomposition reactions are considered the exact reverse of combination reactions due to their structural mechanisms:

1. In a Combination Reaction: Two or more reactants combine to form a single product ($A + B \to AB$):
$$2\text{Mg (s)} + \text{O}_2\text{ (g)} \longrightarrow 2\text{MgO (s)}$$
$$\text{CaO (s)} + \text{H}_2\text{O (l)} \longrightarrow \text{Ca(OH)}_2\text{ (aq)}$$

2. In a Decomposition Reaction: Conversely, a single reactant breaks down into two or more simpler products ($AB \to A + B$):
$$2\text{FeSO}_4\text{ (s)} \xrightarrow{\Delta} \text{Fe}_2\text{O}_3\text{ (s)} + \text{SO}_2\text{ (g)} + \text{SO}_3\text{ (g)}$$
$$\text{CaCO}_3\text{ (s)} \xrightarrow{\Delta} \text{CaO (s)} + \text{CO}_2\text{ (g)}$$

While combination reactions involve bond formation (releasing energy), decomposition reactions involve bond breaking (requiring energy input). Hence, they are polar opposites.

The three modes of decomposition with representative chemical equations are:

1. Thermal Decomposition (Energy supplied as Heat Δ):
$$2\text{Pb(NO}_3)_2\text{ (s)} \xrightarrow{\Delta} 2\text{PbO (s)} + 4\text{NO}_2\text{ (g)} \uparrow + \text{O}_2\text{ (g)} \uparrow$$
(Observation: Emission of brown fumes of nitrogen dioxide gas.)

2. Photolytic Decomposition (Energy supplied as Light):
$$2\text{AgBr (s)} \xrightarrow{\text{Sunlight}} 2\text{Ag (s)} + \text{Br}_2\text{ (g)}$$
(Used in black and white photography; pale yellow silver bromide turns grey.)

3. Electrolytic Decomposition (Energy supplied as Electricity):
$$2\text{H}_2\text{O (l)} \xrightarrow{\text{Electric Current}} 2\text{H}_2\text{ (g)} \uparrow + \text{O}_2\text{ (g)} \uparrow$$
(Hydrogen is liberated at the cathode and oxygen at the anode in a 2:1 volume ratio.)

The key differences are outlined in the comparison table below:

Copper (Cu) is placed above silver (Ag) in the metal activity series, making it more reactive.

When copper turnings or strips are introduced into a colourless silver nitrate solution, copper displaces silver, forming blue-coloured copper(II) nitrate solution and depositing pure, lustrous metallic silver:

$$\text{Cu (s)} + 2\text{AgNO}_3\text{ (aq) [Colourless]} \longrightarrow \text{Cu(NO}_3)_2\text{ (aq) [Blue]} + 2\text{Ag (s)} \downarrow \text{ [Pure Silver]}$$

Definition of Precipitation Reaction:
Any chemical reaction between two clear, aqueous salt solutions that results in the formation of an insoluble solid substance that separates from the liquid solution is called a precipitation reaction. The insoluble solid formed is termed a precipitate.

• Example 1 (Formation of White Barium Sulphate Precipitate):
When aqueous sodium sulphate is mixed with aqueous barium chloride, insoluble white barium sulphate precipitates immediately:
$$\text{Na}_2\text{SO}_4\text{ (aq)} + \text{BaCl}_2\text{ (aq)} \longrightarrow \text{BaSO}_4\text{ (s)} \downarrow \text{ [White ppt]} + 2\text{NaCl (aq)}$$

• Example 2 (Formation of Yellow Lead Iodide Precipitate):
Mixing colourless lead(II) nitrate and potassium iodide solutions produces an insoluble bright yellow precipitate of lead iodide:
$$\text{Pb(NO}_3)_2\text{ (aq)} + 2\text{KI (aq)} \longrightarrow \text{PbI}_2\text{ (s)} \downarrow \text{ [Yellow ppt]} + 2\text{KNO}_3\text{ (aq)}$$

(a) Oxidation:
The chemical process in which a substance gains oxygen or loses hydrogen during a chemical reaction is defined as oxidation.
• Example 1: Oxidation of copper powder on heating in air:
$$2\text{Cu (s)} + \text{O}_2\text{ (g)} \xrightarrow{\Delta} 2\text{CuO (s)}$$
(Copper gains oxygen to form black copper oxide.)
• Example 2: Oxidation of carbon during burning:
$$\text{C (s)} + \text{O}_2\text{ (g)} \longrightarrow \text{CO}_2\text{ (g)}$$
(Carbon gains oxygen to form carbon dioxide.)

(b) Reduction:
The chemical process in which a substance loses oxygen or gains hydrogen during a chemical reaction is defined as reduction.
• Example 1: Reduction of copper oxide by hydrogen gas:
$$\text{CuO (s)} + \text{H}_2\text{ (g)} \xrightarrow{\Delta} \text{Cu (s)} + \text{H}_2\text{O (l)}$$
(Copper oxide loses oxygen to form metallic copper.)
• Example 2: Reduction of zinc oxide by carbon:
$$\text{ZnO (s)} + \text{C (s)} \longrightarrow \text{Zn (s)} + \text{CO (g)}$$
(Zinc oxide loses oxygen to form zinc metal.)

(i) Identification of Element ‘X’ and Black Compound:
• The shiny brown element ‘X’ is Copper metal (Cu).
• The black coloured compound formed is Copper(II) oxide (CuO).

(ii) Balanced Chemical Equation:
When copper powder is heated in a china dish exposed to air, copper combines with oxygen to form a black surface coating of copper(II) oxide:
$$2\text{Cu (s) [Shiny Brown]} + \text{O}_2\text{ (g)} \xrightarrow{\text{Heat (}\Delta\text{)}} 2\text{CuO (s) [Black Coating]}$$
(Note: If hydrogen gas is passed over this hot black CuO, it reduces back to brown copper: $\text{CuO} + \text{H}_2 \to \text{Cu} + \text{H}_2\text{O}$.)

Iron articles rust when exposed to moisture (water vapour) and oxygen present in the atmosphere, forming hydrated iron(III) oxide ($\text{Fe}_2\text{O}_3 \cdot x\text{H}_2\text{O}$), a flaky, porous substance that causes iron to corrode and deteriorate over time.

Reason for Applying Paint:
1. Paint provides an impermeable protective surface barrier between the bare iron metal and atmospheric oxygen and moisture.
2. Without direct contact with air and moisture, electrochemical oxidation cannot initiate.
3. Consequently, corrosion is prevented, conserving the structural integrity, durability, and appearance of iron articles.

Food products containing edible oils and fats (such as potato chips, wafers, roasted nuts) undergo spontaneous oxidation when exposed to atmospheric oxygen over time.

This oxidation leads to the decomposition of fats into foul-smelling, volatile fatty acids, causing unpleasant changes in both taste and odor. This condition is called rancidity, rendering the food unpalatable and unhealthy.

Why Nitrogen Gas is Used:
1. Nitrogen ($\text{N}_2$) is an unreactive, inert gas at room temperature.
2. Flushing food packages with nitrogen displaces atmospheric oxygen inside the airtight pouch.
3. In the complete absence of oxygen, oxidative degradation is halted, preserving the crispiness, aroma, and freshness of packaged snacks for several months.

(a) Corrosion:
The gradual deterioration and degradation of refined metals caused by spontaneous chemical or electrochemical reactions with environmental agents (such as moisture, atmospheric gases, and acids) is known as corrosion.
• Examples:
1. Rusting of Iron: Formation of reddish-brown hydrated ferric oxide on iron ($4\text{Fe} + 3\text{O}_2 + 2x\text{H}_2\text{O} \to 2\text{Fe}_2\text{O}_3 \cdot x\text{H}_2\text{O}$).
2. Tarnishing of Silver: Formation of a black coating of silver sulphide ($\text{Ag}_2\text{S}$) upon contact with airborne traces of hydrogen sulphide.
3. Green Coating on Copper: Formation of basic copper carbonate ($\text{CuCO}_3 \cdot \text{Cu(OH)}_2$) when copper interacts with moist carbon dioxide.

(b) Rancidity:
The natural oxidation of fats and oils present in food items when exposed to atmospheric air, resulting in unpleasant odors, rancid flavors, and loss of nutritional value, is called rancidity.
• Example: Cooking oil or butter turning sour, smelling rancid, and tasting stale after prolonged storage in unsealed containers.
• Prevention: Adding antioxidants (BHA, BHT), vacuum packing, flushing airtight packets with nitrogen gas, and refrigerating foodstuffs.